A galvanic cell (also called a voltaic cell) is a device that converts chemical energy into electrical energy through a spontaneous redox reaction. It is one of the core ideas in IB Chemistry Topic 9 (Redox Processes) and appears frequently in Paper 2 and 3. Galvanic cells help explain how batteries work, why electrons flow, and how electrical potential is generated from chemical reactions. Understanding galvanic cells strengthens your grasp of electron transfer, oxidation, reduction, and electrode behavior.
What Is a Galvanic Cell?
A galvanic cell is an electrochemical cell that generates electrical energy from a spontaneous redox reaction.
Key features:
- Chemical energy → electrical energy
- Electrons flow through an external circuit
- Two half-cells are required
- A salt bridge maintains charge balance
Galvanic cells power everyday batteries, making the concept both practical and essential.
The Structure of a Galvanic Cell
A classic IB-style galvanic cell has four main components:
1. Two half-cells
Each half-cell contains:
- A metal electrode
- An aqueous solution of its ions
Examples:
- Zn(s) in Zn²⁺(aq)
- Cu(s) in Cu²⁺(aq)
2. Electrodes
- Anode: where oxidation occurs
- Cathode: where reduction occurs
Mnemonic: OIL RIG
Oxidation Is Loss (electrons)
Reduction Is Gain (electrons)
3. External wire
Allows electrons to flow from anode → cathode.
4. Salt bridge
A tube filled with an electrolyte that:
- Maintains electrical neutrality
- Allows ions to flow
- Completes the circuit
Without a salt bridge, the reaction stops.
How a Galvanic Cell Works
Step 1: Oxidation at the anode
The metal loses electrons.
Example:
Zn(s) → Zn²⁺(aq) + 2e⁻
Step 2: Electrons travel through the wire
Electrons move from the anode to the cathode.
This electron flow generates electricity.
Step 3: Reduction at the cathode
Metal ions gain electrons.
Example:
Cu²⁺(aq) + 2e⁻ → Cu(s)
Step 4: Salt bridge maintains charge balance
- Anions move to the anode compartment
- Cations move to the cathode compartment
This prevents charge buildup.
Example: The Daniell Cell (Zn–Cu Cell)
A common IB example uses zinc and copper electrodes:
Anode (oxidation):
Zn(s) → Zn²⁺(aq) + 2e⁻
Cathode (reduction):
Cu²⁺(aq) + 2e⁻ → Cu(s)
Electron Flow:
Zn → Cu
Ion Flow:
- Anions → anode
- Cations → cathode
Result:
Copper plates onto the cathode, zinc dissolves at the anode, and electrical energy is produced.
Cell Potential (Voltage)
The voltage of a galvanic cell depends on the difference between the standard reduction potentials of the two half-reactions:
E°cell = E°cathode − E°anode
- A positive value indicates a spontaneous reaction.
- Higher voltage means a stronger driving force for electron flow.
IB exams often ask students to calculate E°cell from given reduction potentials.
Why Galvanic Cells Are Important
1. Foundation of batteries
All real batteries operate as galvanic cells.
2. Shows how chemical reactions create electricity
Spontaneous reactions drive electron flow.
3. Teaches redox chemistry
Electrodes clearly show oxidation and reduction processes.
4. Helps understand equilibrium
As the cell runs, concentrations change and voltage decreases.
5. Appears in many exam questions
IB loves diagram-based galvanic cell problems.
Common IB Misunderstandings
“Electrons flow through the salt bridge.”
False—ions flow through the salt bridge, not electrons.
“Cathode is always positive.”
This is true for galvanic cells, but not for electrolytic cells.
“Oxidation happens at the cathode.”
Oxidation always occurs at the anode.
“The reaction continues indefinitely.”
It stops when reactants run out or concentrations change too much.
FAQs
Why is the anode negative in a galvanic cell?
Because it pushes electrons away as oxidation releases electrons.
Does the cathode gain mass?
Yes—metal ions are reduced onto its surface.
Why is a salt bridge necessary?
To prevent charge buildup and allow the reaction to continue.
Conclusion
A galvanic cell converts chemical energy into electrical energy using a spontaneous redox reaction. With an anode for oxidation, a cathode for reduction, and a salt bridge maintaining balance, galvanic cells demonstrate the core principles of redox chemistry. Mastering this concept helps IB Chemistry students understand battery technology, electron flow, and electrochemical behavior.
